Section Periodic Trends Objective: - Compare period & group trends for shielding, atomic radius, ionic radius, ionization energy, & electronegativity. Shielding (or screening): The valence e- are blocked from the full positive charge of the nucleus (effective nuclear charge) by the inner (core) e-. As the average number of core e- increases, the effective nuclear charge decreases. –Concept of shielding will play a large role in a lot of the trends.

Trend within the period (left to right): Generally decreases Why: –Number of energy levels & core e- stays the same, but nucleus is increasing. –Increased attraction between nucleus and valence e-. Trend down a group: Generally increases Why: –Number of energy levels & core e- increases. –Valence e- farther from nucleus & more blocked by inner e-. Shielding:

Atomic Radius: ½ the distance between adjacent nuclei of identical atoms. Trend within the period (left to right): Generally decreases Why: –Number of energy levels & core e- stays the same, but nucleus is increasing. –Increased attraction between nucleus and valence e-. –This attraction pulls the e- closer to the nucleus and makes the atom smaller. Trend down a group: Increases Why: –Number of energy levels increases & core e- increases. –Each energy level is larger than the next. –Valence e- farther from the nucleus and more blocked by the inner e-.

Examples – Place each group of elements in order of increasing atomic radius: 1.S, Al, Cl, Mg, Ar, Na 2.K, Li, Cs, Na, H 3.Ca, As, F, Rb, O, K, S, Ga

Examples – Place each group of elements in order of increasing atomic radius: 1.S, Al, Cl, Mg, Ar, Na Ar < Cl < S < Al < Mg < Na 2.K, Li, Cs, Na, H H < Li < Na < K < Cs 3.Ca, F, As, Rb, O, K, S, Ga F < O < S < As < Ga < Ca < K < Rb

Ionic Radius: – Distance between the nucleus and the outermost electron in ions (can’t be determined directly). Trend between atom & ion: –Cations are smaller than original atom. (Losing e-, the atom has unequal positive charge that attracts the valence e- closer to the nucleus.) –Anions are larger than original atom and cations. (Adding negative e-, adds to repulsion between valence e-, pushing them apart.) Trend within the period (left to right): Representative Elements → Decreases. –Cations: size decreases. –Anions: the size drastically increases compared to the positive ions, and then decreases across the period. Trend down a group: Increases for both cations & anions. –Same reason as atomic radii trend.

Write electron configurations for the following ions: 1.Al 3+ 2.S 2- 3.Li + 4.Br - 5.Fe 2+ 6.Fe 3+

Write electron configurations for the following ions: 1.Al 3+ 1s 2 2s 2 2p 6 2.S 2- [Ne]3s 2 3p 6 3.Li + 1s 2 4.Br - [Ar]4s 2 3d 10 4p 6 5.Fe 2+ [Ar]3d 6 6.Fe 3+ [Ar]3d 5

Examples – Choose the larger species in each case: 1.Na or Na + 2.Br or Br - 3.N or N 3- 4.O - or O 2- 5.Mg 2+ or Sr 2+ 6.Mg 2+ or O 2- 7.Fe 2+ or Fe 3+

Examples – Choose the larger species in each case: 1.Na or Na + 2.Br or Br - 3.N or N 3- 4.O - or O 2- 5.Mg 2+ or Sr 2+ 6.Mg 2+ or O 2- 7.Fe 2+ or Fe 3+

Ionization Energy: Energy required to remove an electron from a gaseous atom (also called First Ionization Energy, I 1 ) Na (g) kJ  Na + (g) + e - The second ionization energy, I 2, is the energy required to remove the next available electron: Na + (g) kJ  Na 2+ (g) + e - NOTICE: –Ionization Energy increases for each electron removed from the same element. –The larger ionization energy, the more difficult it is to remove the electron.

Variations in Successive Ionization Energies There is a sharp increase in ionization energy when a core electron is removed. Notice the large increase after the last valence electron is removed. This chart can be used to determine the number of valence electrons in an atom of an element.

Trend within the period: Increases Why: –Electrons are more difficult to remove from smaller atoms. –Closer to the nucleus and increased nuclear charge. Trend down a group: Decreases Why: –Electrons are easier to remove from large atoms. –Farther away from the nucleus so less energy is needed to remove them. Notice the trend in ionization energy is inversely related to trends in atomic radii.

Examples – Put each set in order of increasing first ionization energy: 1.P, Cl, Al, Na, S, Mg 2.Ca, Be, Ba, Mg, Sr 3.Ca, F, As, Rb, O, K, S, Ga

Examples – Put each set in order of increasing first ionization energy: 1.P, Cl, Al, Na, S, Mg 2.Ca, Be, Ba, Mg, Sr 3.Ca, F, As, Rb, O, K, S, Ga 1. Na < Al < Mg < S < P < Cl 2.Ba < Sr < Ca < Mg < Be 3.Rb < K < Ca < Ga < As < S < O < F

ELECTRONEGATIVITY: Ability of an atom to attract electrons in a chemical bond to itself. Chemist Linus Pauling set electronegativities on a scale. –0.7 (Cs) to 4.0 (F) –Used to help determine types of bonding (ionic or covalent) that are occurring in a compound. –Noble gases are not usually given electronegativity values. Trend within the period: Increases Why: –Atoms become smaller, so shared electrons are closer to the nucleus. Trend down a group: Decreases Why: –Atoms become larger, so shared electrons are farther from the nucleus.

Electronegativity

Examples – put each set in order by increasing electronegativity: 1.Na, Li, Rb, K, Fr 2.Cl, Ca, F, P, Mg, S, K

Examples – put each set in order by increasing electronegativity: 1.Na, Li, Rb, K, Fr 2.Cl, Ca, F, P, Mg, S, K 1.Fr < Rb < K < Na < Li 2.K < Ca < Mg < P < S < Cl < F

Review: 1.As you move across a period, left to right, describe what generally happens (decreases, increases, or remains the same) to: a.Number of valence electrons b.Ionization energy c.Atomic radius 2. Give a brief explanation for your answers to a-c.

3. Identify the element from the clues given: a.This element has a smaller atomic radius than phosphorous, it has a smaller ionization energy than fluorine, and is chemically similar to iodine. b.This element has the smallest ionization energy of any element in Period 4.